1. The modern periodic law; groups and periods
Modern periodic law: the properties of elements are a periodic function of their atomic numbers. Atomic number fixes the number of electrons, hence the valence electrons, which decide properties (argon, 39.9, correctly comes before potassium, 39.1).
- Group (18 vertical columns): same number of valence electrons; similar properties. Numbered 1–18 (IUPAC) or IA–VIIA, 0.
- Period (7 horizontal rows): same number of shells. Period 1: 2 elements; periods 2 and 3: 8 each; period 4 starts K, Ca.
- From configuration: number of shells = period; valence electrons = group (1, 2) or 10 + valence electrons (groups 13–18).
| Element | Configuration | Period | Group |
|---|---|---|---|
| N (7) | 2, 5 | 2 | 15 |
| Mg (12) | 2, 8, 2 | 3 | 2 |
| Ca (20) | 2, 8, 8, 2 | 4 | 2 |
Key terms
- Modern periodic law
The properties of elements are a periodic function of their atomic numbers.
- Period
A horizontal row of the periodic table; its elements have the same number of electron shells.
2. Nuclear charge and number of shells
- Across a period: shells stay the same; nuclear charge increases by one each step → valence electrons pulled in more strongly.
- Down a group: a new shell is added each step → valence electrons farther away and shielded by inner shells → weaker pull, despite the larger nuclear charge.
Every trend in this chapter follows from these two statements.
| Direction | Pull on valence electrons |
|---|---|
| Across a period | Stronger |
| Down a group | Weaker |
Key terms
- Nuclear charge
The positive charge of the nucleus, equal to the number of protons (atomic number).
- Shielding
The screening of outer electrons from the nucleus by electrons in inner shells.
3. Atomic size
Atomic radius: distance from the centre of the nucleus to the outermost shell. Decreases across a period (Na 186 pm → Cl 99 pm); increases down a group (Li < Na < K). Noble-gas radii are van der Waals radii and are not compared with the bonded radii of the other elements in a period.
Cations are smaller than their atoms (Na⁺ < Na); anions are larger (Cl⁻ > Cl).
| Pair | Larger | Reason |
|---|---|---|
| Na, Cl | Na | Smaller nuclear charge, same shells |
| Li, K | K | More shells |
Key terms
- Atomic radius
The distance from the centre of the nucleus to the outermost shell of an atom.
4. Metallic and non-metallic character
Metallic character: tendency to lose electrons (form cations). Non-metallic character: tendency to gain electrons (form anions).
- Across a period: metallic ↓, non-metallic ↑ (period 3: Na, Mg, Al metals; Si metalloid; P, S, Cl non-metals).
- Down a group: metallic ↑, non-metallic ↓.
- Oxides change from basic (Na₂O, MgO) through amphoteric (Al₂O₃) to acidic (SO₂, Cl₂O₇) across period 3.
| Element | Character | Oxide |
|---|---|---|
| Na | Metal | Basic |
| Al | Metal | Amphoteric |
| S | Non-metal | Acidic |
Key terms
- Metallic character
The tendency of an atom to lose electrons and form a positive ion.
5. Ionisation potential, electron affinity and electronegativity
| Property | Definition | Across a period | Down a group |
|---|---|---|---|
| Ionisation potential | Energy needed to remove an electron from the outer shell of an isolated gaseous atom | Increases | Decreases |
| Electron affinity | Energy released when an isolated gaseous atom gains an electron | Increases | Decreases (generally) |
| Electronegativity | Tendency of an atom in a molecule to attract the shared pair to itself | Increases | Decreases |
Noble gases: highest ionisation potential in each period, almost zero electron affinity. Fluorine: most electronegative element. Chlorine’s electron affinity is slightly higher than fluorine’s.
| Highest in period 3 | Element |
|---|---|
| Ionisation potential | Ar |
| Electronegativity | Cl |
| Atomic size (excluding Ar) | Na |
Key terms
- Ionisation potential
The energy required to remove an electron from the outermost shell of an isolated gaseous atom.
- Electron affinity
The energy released when an isolated gaseous atom gains an electron to form a negative ion.
- Electronegativity
The tendency of an atom in a molecule to attract the shared pair of electrons towards itself.
6. Alkali metals and halogens
Group 1 (Li, Na, K): one valence electron; size ↑ and ionisation potential ↓ down the group, so reactivity increases (Li < Na < K). With water: 2M + 2H₂O → 2MOH + H₂.
Group 17 (F, Cl, Br, I): seven valence electrons; size ↑ down the group, so attraction for an extra electron ↓ and reactivity decreases (F > Cl > Br > I). Displacement: Cl₂ + 2KBr → 2KCl + Br₂; Br₂ + 2KI → 2KBr + I₂.
| Family | Trend down the group |
|---|---|
| Alkali metals | More reactive |
| Halogens | Less reactive |
Key terms
- Displacement (halogens)
A more reactive halogen removes a less reactive halogen from a solution of its salt.
Exam tips
Common mistakes
- Saying atomic size increases across a period.
- Mixing up electron affinity (energy released, isolated atom) and electronegativity (attraction in a molecule).
- Comparing argon’s radius with chlorine’s.
- Saying halogens become more reactive down the group.
How marks are usually awarded
- Trend questions (1 mark each): the correct order plus a reason using nuclear charge or shells.
- Definitions (1 mark): include “isolated gaseous atom” where needed.
- Placing an element (2 marks): configuration, then period and group.